Solutions State of State of State of Example Solution Solute Solvent Air, natural gas Gas Gas Gas Rubbing alcohol, antifreeze Liquid Liquid Liquid Brass Solid Solid Solid Carbonated water (soda) Liquid Gas Liquid Seawater, sugar solution Liquid Solid Liquid Hydrogen in platinum Solid Gas Solid
Components of Solution Relationships Between Amounts of Solute, Solvent and Solution Molar Mass Density
nsolute msolute Vsolute Molar Mass Densitynsolvent msolvent Vsolventnsoln msoln Vsoln molessolute Molarity (M ) volum esolutionin L
nsolute msolute Vsolute Molar Mass Densitynsolvent msolvent Vsolventnsoln msoln Vsoln m oles solute Molality (m) m ass solvent(kg)
nsolute msolute Vsolute Molar Mass Densitynsolvent msolvent Vsolventnsoln msoln Vsoln m ass A m ass A m ass % A 100 100 m ass A m ass B m ass C total m ass solution
nsolute msolute Vsolute Molar Mass Densitynsolvent msolvent Vsolventnsoln msoln Vsoln m oles A m oles A Mole fraction A m oles solution m oles A m oles B m olesC
1. Separating the solute into its individual components (expanding the solute).2. Overcoming intermolecular forces in the solvent to make room for the solute (expanding the solvent).3. Allowing the solute and solvent to interact to form the solution.
Steps 1 and 2 require energy, since forces must be overcome to expand the solute and solvent.Step 3 usually releases energy.Steps 1 and 2 are endothermic, and step 3 is often exothermic.
Explain why water and oil (a long chain hydrocarbon) do notmix. In your explanation, be sure to address how ΔH plays arole.
Energy Terms of Various Typesof Solutes and Solvent H1 H2 H3 Hsoln OutcomePolar solute, polar Large Large Large, Small Solutionsolvent negative formsNonpolar solute, polar Small Large Small Large, No solutionsolvent positive formsNonpolar solute, Small Small Small Small Solutionnonpolar solvent formsPolar solute, nonpolar Large Small Small Large, No solutionsolvent positive forms
In GeneralOne factor that favors a process is an increase in probability of the state when the solute and solvent are mixed.Processes that require large amounts of energy tend not to occur.Overall, remember that “like dissolves like”.
Henry’s law: C = kP C = concentration of dissolved gas k = constant P = partial pressure of gas solute above the solutionAmount of gas dissolved in a solution is directly proportional to the pressure of the gas above the solution.
Temperature Effects (forAqueous Solutions)• Although the solubility of most solids in water increases with temperature, the solubilities of some substances decrease with increasing temperature.• Predicting temperature dependence of solubility is very difficult• Solubility of a gas in solvent typically decreases with increasing temperature.
Temperature Dependence ofSolubility of Solids in AqueousSolutions
Temperature Dependence ofSolubility of Gases in AqueousSolutions
Temperature Dependence ofSolubility in Aqueous Solutions Although the solubility of most solids in water increases with temperature, the solubilities of some substances decrease with increasing temperature. Predicting temperature dependence of solubility is very difficult. Solubility of a gas in solvent typically decreases with increasing temperature.
Colligative Properties Properties that depend ONLY on the number and NOT in the nature of solute particles. Vapor Pressure Lowering Boiling Point Elevation Freezing Point Elevation Osmotic Pressure
Equilibrium Vapor Pressure • The pressure of the vapor due to the evaporation of the liquid (solid) above the same liquid (or solid) in a closed container. • The rate of evaporation is equal to the rate of condensation • Is it affected by • Strength of intermolecular forces in the liquid? • The stronger the IMFA, the lower the VP • The weaker the IMFA, the higher the VP • Temperature? • Higher temperature increases energy to overcome IMFA, higher VP • Surface area of the container? • VP is force/area • The greater the surface area, more molecules evaporate, the force increases proportionally • No effect
Vapor Pressure Lowering + Non Pure Volatile Solution Solvent Solute
Vapor Pressure Lowering • Evaporation occurs at the surface at any temperature • Some of the non-volatile solute particles take the place of the solvent particles at the surface • Non-volatile solute particles increase the intermolecular forces of attraction between solute and solvent • Decreases the number of solvent particles evaporating • Lower Vapor pressure for the solution
Raoult’s Law• As the mole fraction of the solvent increases, the vapor pressure of the solution approaches that of the pure solvent.• Dilute solutions have properties closer to the solvent
Raoult’s Law • Nonvolatile solute • Volatile solute • Volatile solvent • Volatile solvent • No need to indicate solvent/solute
Ideal vs. Non-ideal Solutions • Does not follow • Does not follow Raoult’s Law Raoult’s Law• Follows Raoult’s Law • Weak solute-solvent • Strong solute-solvent• Non-polar/non-polar interactions interactions• Similar shape and size • Both want to stay in • Both want to stay in the vapor phase the liquid phase • Predicted Ptotal is • Predicted Ptotal is lower higher • Negatively deviating • Positively Deviating • Predominantly polar • Predominantly non- interactions polar interactions
Summary of Solute-SolventInteractions in Solution Deviation Interactive Forces T for fromBetween Solute (A) and Hsoln Solution Example Raoult’s Solvent (B) Particles Formation Law None Benzene-A A, B B A B Zero Zero (ideal toluene solution) NegativeA A, B B<A B Positive Negative Acetone-water (exothermic) Positive Ethanol-A A, B B>A B Negative Positive (endothermic) hexane
Boiling Point It is the temperature when the vapor pressure of the liquid is equal to the prevailing atmospheric pressure Is it affected by ◦ IMFA in liquid? Stronger IMFA, required more energy to increase VP, higher BP ◦ Prevailing atmospheric pressure? Higher atmospheric pressure, higher BP ◦ Ground elevation? Higher ground elevation, lower atmospheric pressure, lower BP ◦ Prevailing temperature? Substance must reach a certain temperature before it boils Depends on the three previous factors
Boiling Point Elevation + Non Volatile Pure Solvent /Nonelectrolyte Solution Solute
Boiling Point Elevation Nonvolatile solute elevates the boiling point of the solvent. ΔTb = Kbmsolute ΔTb = boiling-point elevation Kb = molal boiling-point elevation constant m = molality Boiling Solvent Normal of solution o K , C/m b Point, oC Water 100.0 0.512 Acetic acid (CH3COOH) 118.9 3.1 Benzene (C6H6) 80.1 2.53 Chloroform (CHCl3) 61.2 3.63
Freezing Point Depression When a solute is dissolved in a solvent, the freezing point of the solution is lower than that of the pure solvent. ΔTf = Kfmsolute ΔTf = freezing-point depression Kf = molal freezing-point depression constant m = molality of solute Normal Freezing Solvent Kf, oC/m Point, oC Water 0 1.86 Ethyl ether (C4H10O) -116.2 1.79 Benzene (C6H6) 5.5 5.12 Chloroform (CHCl3) -63.5 4.70
Osmotic PressureOsmosis – flow of solvent into the solution through a semipermeable membrane. = MRT = osmotic pressure (atm) M = molarity of the solution R = gas law constant T = temperature (Kelvin)
Colligative Properties of Electrolytes van’t Hoff Factor, i The relationship between the moles of solute dissolved and the moles of particles in solution is usually expressed as: moles of particles in solution i = moles of solute dissolved
Colligative Properties ofElectrolytes Electrolyte i, (expected) i, (observed) NaCl 1.9 MgCl2 2.7 MgSO4 1.3 i FeCl3 3.4 HCl 1.9 Glucose* 1.0 Tb iK b m Tf iK f m iMRT
Ion Pairing At a given instant a small percentage of the sodium and chloride ions are paired and thus count• as a pairing is most important in concentrated solutions. Ion single particle.• As the solution becomes more dilute, the ions are farther apart and less ion pairing occurs.• Ion pairing occurs to some extent in all electrolyte solutions.• Ion pairing is most important for highly charged ions.
Colloids• A suspension of tiny particles in some medium.• Tyndall effect – scattering of light by particles.• Suspended particles are single large molecules or aggregates of molecules or ions ranging in size from 1 to 1000 nm.
ColloidsCoagulation Destruction of a colloid. Usually accomplished either by heating or by adding an electrolyte.